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Chemistry > 3.1 > Flashcards

3.1 Flashcards

1
Q

Problems with Antoine Laurent de Lavoisier

A

Included light and heat, thought to be material

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2
Q

Jons Jakob Berzelius

A

Table of atomic weights

Letter based symbols

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3
Q

Johan Wolfgang debereiner

A

Noticed groups of 3 elements ordered by atomic weight, middle had weight and properties, average of other 2

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4
Q

John Newlands

A

Table, arranged in order of relative atomic weight

Element have similar properties those 8 places in front law of octaves

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5
Q

Mendeleev

A

Modern based on
Elements ordered by atomic weights
Periodically
Columns, similar properties, groups
Gaps left where none fitted, predicted properties of undiscovered elements
Gallium, germanium, scandium matched these predictions
Order reversed when properties didn’t fit

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6
Q

Henry Moseley

A

Determined atomic number for all known elements
Properties vary periodically with atomic number rather than weight.
Fixed elements Mendeleev had switched around

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7
Q

Glenn Seaborg

A

Discovered transuranic elements from plutonium 94 to nobelium 102.
Placed actinide series below lanthanides series at bottom of table.

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8
Q

List makers of period table

A 
Antoine Laurent de Lavoisier 
Jons Jakob Berzelius 
Johann Wolfgang dobereiner 
John Newlands
Mendeleev
Henry Moseley
Glenn Seaborg
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9
Q

What is the table ordered by?

A

Increasing atomic number starts with hydrogen

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10
Q

Describe periodicity

A

Horizontal row
Gradual changes in properties across a period in the table
Repeated across each period
Repeated pattern= periodicity
Repeating properties across periods
Predict properties metals to non metals across

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11
Q

Describe groups of the table

A

Vertical columns

Elements with similar properties

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12
Q

Describe semi metals, metalloids

A

Non-metals
in between metals and non metals boron etc
Properties between those of metals and non metals

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13
Q

Describe atomic radius

A

Atomic radii increase too to bottom of group number of shells increases down group and nuclear charge
This decreases the size of individual shells

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14
Q

Describe electronegativity

A

Increases across period
Decreases down group
Nuclear charge increases electronegativity increases,
Distance from nucleus increases low electronegativity electronpositive

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15
Q

Equation for 1st ionisation energy

A

X —-> x+ + e-

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16
Q

Equation of second ionisation energy

A

X+ ——-> x2+ + e-

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17
Q

Equation for third ionisation energy

A

X2+ ——> x3+ + e-

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18
Q

What state is always used for ionisation energies?

A

Gas

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19
Q

Define enthalpy

A

Energy change per mole

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20
Q

Define ionisation energy

A

Energy to remove electron from an atom

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21
Q

Define first ionisation energy

A

The amount of energy required to remove one electron from each atom in a mole of atoms; elements in gaseous state

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22
Q

Describe the trend in first ionisation energy

A

Easier from top to bottom of group
More shells less attraction between outer electron and positive nucleus

Harder across more protons higher force of attraction

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23
Q

With successive ionisation energies when will big differences in energy occur?

A

Between electron shells because the shell closer to the nucleus will have a greater attraction for the electrons and less shielding from complete shells.

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24
Q

3 factors ionisation energy trends can be explained by

A

Distance from the nucleus
The nuclear charge/ attraction
The amount of shielding

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25
Q

What did Antoine-Laurent de Lavoisier do?

A

Produced first modern chemical textbook
Extensive list of elements ‘substances that cannot be broken down further’
Distinguished between metals and non metals

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26
Q

What happens to first ionisation energy down a group and why?

A

Decreases, electron is further from nucleus, more shielding, even though there is an increase in nuclear charge.

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27
Q

What happens to first ionisation energy across a period?

e.g. Li to Ne

A

Overall increase in ionisation energy due to the increase in nuclear charge but the same distance from the nucleus.
Drop from Be to B due to shielding from full 2s orbital
Drop from N to O due to electron repulsion when P electrons pair up, easier to remove.

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28
Q

Why do giant metallic lattices have high melting and boiling points?

A

Electrons are free to move throughout the structure but positive ions remain where they are.

The attraction between positive ions and negative delocalised electrons is very strong.
A high temperature is needed to overcome the metallic bonds and dislodge the ions from their rigid positions in the lattice.

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29
Q

Why are giant metallic lattices good electrical conductors?

A

The delocalised electrons can mover freely throughout the lattice.
This allows it to conduct even in solid state.

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30
Q

Why are giant metallic lattices both malleable and ductile?

A

Delocalised electrons, can move the structure has a degree of give which allows atoms or layers to slide past each other.

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31
Q

Define ductile

A

Can be drawn out of stretched

Permits metals to be drawn into wires.

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32
Q

Define malleable

A

Can be hammered into different shapes

Many metals can be pressed into different shapes or hammered into thin sheets.

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33
Q

What happens as you move across periods 2 and 3?

A

Metal element to non metal elements

Solid to gases

34
Q

How does melting point change across periods 2 and 3?

A

Group 1-14 melting points increase steadily, because they have giant structures. If the structure is metallic as the nuclear charge increases so does the number of electrons in the outer shell, stronger attraction.
Giant covalent lattice, more electrons to form covalent bonds.

Group 14-15 sharp decrease in melting point, simple molecular structures, weak intermolecular forces.

Group 15-18 melting points remain low, simple molecular structures.

35
Q

Describe the structure of graphite:

A

2D giant lattice, one carbon atom thick of interlocking hexagonal carbon rings. Extremely strong, very light and can conduct electricity.

36
Q

Key properties of group 2 metals

A

Reasonably high melting and boiling points
Light metals, low densities
Form colourless white compounds

37
Q

Where is the highest energy level electron in group 1 and 2 metals?

A

s-sub shell

They are in the S block of the periodic table

38
Q

Difference between the electron configuration of group two metals and that of a noble gas:

A

2 electrons more

An outer shell containing 2 electrons

39
Q

What must happen to group 2 elements for them to have the same electron configuration as a noble gas?

A

Lose 2 electrons.

40
Q

Why does reactivity increase down group 2?

A

Each successive element has its outer shell electrons in a higher energy level, has a larger atomic radius and feels more shielding from the positive nucleus.

41
Q

What are group 2 metals oxidised to form?

A

2+ ions

42
Q

How do group 2 elements react with oxygen?

A

Vigorously
Redox reaction
Product = ionic oxide
e.g. 2Ca + O2 = 2CaO

43
Q

How do group 2 elements react with water?

A

All except beryllium react to form hydroxides with the general formula M(OH)2. Hydrogen gas is formed.

Mg reacts slowly, down the group they react more vigorously.

Redox reaction, metal is oxidised and one hydrogen per water molecule is reduced.

44
Q

How do group 2 metals react with dilute acids?

A

All except Be react to form a salt and hydrogen gas.

More vigorous down group.

45
Q

What does Calcium and hydrochloric acid form?

A

Calcium chloride and hydrogen

46
Q

How do group 2 oxides react with water?

A

Form metal hydroxides

Soluble in water because they release OH- ions. Typical pH is 10-12.

47
Q

How does the solubility of metal hydroxides of group 2 differ?

A

Increases down the group. When a hydroxide is more soluble it will release more OH- ions, makes a more alkaline solution with a higher ph.
e.g. Beryllium hydroxide is insoluble
Magnesium hydroxide is slightly soluble, solution is dilute with a small OH- concentration.
Barium hydroxide much more soluble, higher OH- concentration, so forms a more alkaline concentration.

48
Q

In what industries are metal hydroxides important?

A

Neutralising acids

Construction industry

49
Q

How are metal hydroxides used to neutralise acidic soil?

A

Calcium hydroxide, used by farmers and gardeners as ‘lime’ to reduce the acidity levels.

50
Q

How are group 2 compounds used as indigestion remedies?

A

Human stomachs contain a small amount of hydrochloric acid. Indigestion = a build up of too much acid.
Many remedies including magnesium hydroxide. Neutralises the excess acid producing a salt and water.

51
Q

What are the building and construction uses of group 2 metal compounds?

A

Calcium carbonate, present in limestone and marble. Calcium carbonate is used in manufacture of glass and steel.

52
Q

What is a draw back of using group 2 metal carbonates in building?

A

Readily react with acids. Most rain water has an acidic pH which leads to a gradual erosion of objects made using limestone or marble such as buildings or statues.

53
Q

Key properties of group 17 elements

A

Low melting and boiling points

Exist as diatomic molecules

54
Q

Describe the trend in boiling point of the halogens

A

Down group increases. Physical state changes from gas to liquid to solid. Each successive element has an extra shell of electrons, higher level of London forces between molecules.

55
Q

What is the last part of the electron configuration of the halogens?

A

P5

56
Q

Describe the reactivity of the halogens

A

Very reactive and highly electronegative.
very good at attracting and capturing electrons. Strong oxidising agents.
During reactions they gain an electron to form a 1- ion and obtain noble gas configuration.
Reactivity and oxidising power decreases down group:
atomic radius increases, nuclear pull is further away; electron shielding increases; ability to gain electron in p sub-shell decreases.

57
Q

What are halide ions?

A

Cl-, Br-, I-

58
Q

What reactions occur between halide ions and halogens?

A

Redox

59
Q

What happens during halogen, halide redox reactions?

A

More reactive halogen oxidises and displaces a halide of a less reactive halogen, displacement reaction.

60
Q

What is an indication a halogen halide displacement reaction has occurred?

A

Halogens form different coloured solutions, so colour changes. Mixture is usually shaken with an organic solvent e.g. cyclohexane to distinguish between bromine and iodine.

61
Q

What colour is Cl2 in water?

A

Pale green

62
Q

What colour is Cl2 in cyclohexane?

A

Pale green

63
Q

What colour is Br2 in water?

A

Orange

64
Q

What colour is Br2 in cyclohexane?

A

Orange

65
Q

What colour is I2 in water?

A

Brown

66
Q

What colour is I2 in cyclohexane?

A

Violet

67
Q

What halide ions does chlorine oxidise?

A

Br- and I-

68
Q

How does Chlorine react with water in water purification?

A

Kills bacteria making it safer to drink
forms Hydrochloric acid and Chloric 1 acid.
Disproportionation as chlorine is both oxidised and reduced.

69
Q

Formula of hydrochloric acid

A

HCl

70
Q

Formula of chloric (I) acid

A

HClO

71
Q

How does chlorine react with cold dilute aqueous sodium hydroxide?

A

Bleach formation
Chlorine is only slightly soluble in water, mild bleaching action. Household bleach is formed when dilute aqueous sodium hydroxide and chlorine react together at room temperature. Chlorine is both oxidised and reduced.

72
Q

Define qualitative tests

A

Tell you which ions are present but not how many

73
Q

Carbonate ion

A

CO32-

74
Q

How would you test for carbonate ions?

A

They react with acids.
Add dilute strong acid to carbonate. Collect gas formed and pass through water.
Fizzing/ colourless gas formed.
Turns limewater cloudy.

75
Q

Equation for carbonate ions reacting with acids

A

CO32- + 2H+ = H2O + CO2

76
Q

How would you test for sulphate ions?

A

Add dilute hydrochloric acid and barium chloride.

A white precipitate of barium chloride is formed.

77
Q

Sulphate ion

A

SO42-

78
Q

Equation of sulphate ions reacting with barium ions

A

Ba2+ + SO42- = BaSO4

79
Q

How would you test for halide ions?

A

Dissolve suspected halide in water.
Add an aqueous solution of solver nitrate.
Note the colour of any precipitate formed.
If it is hard to distinguish add ammonia.
silver chloride = white precipitate, soluble in NH3
Silver bromide = cream precipitate soluble in concentrated NH3.
Silver Iodide = Yellow precipitate, insoluble in dilute and concentrated NH3.

80
Q

How would you test for ammonium ions?

A

Add sodium hydroxide solution to the suspected ammonium compound and warm gently.
test any gas evolved with red litmus.
Ammonium gas would turn red litmus blue, has a distinctive smell, gas is hazardous.

Chemistry (7 decks)

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